1. A
Determine the oxidation number for each atom in the following molecules
H2S
P2O5 S8
SCl2
Na2SO3
6. SO4-2
7. NaH
Cr2O7-2
SnBr4
10. Ba(OH)2
For practice before class

2. 12/05/06 Electrochemistry 19.9-19.13 p 941-955A
Intersection 14
12/05/06
Electrochemistry p

3. December in Studio S M Tu W Th F A 12/5 Exam 3 12/6 Studio
12/8 Polymers; check out
12/11 Poster session, paper due
12/12 final IS
12/13 In-class assignment
12/17 Review session 7-9 pm
12/19 Final exam 8-10 am

4. Watershed Poster Session
Monday, December 11 in USB 2165
Board (4 ft x 4ft), easel, pins
Set up by 1:10 and 3:10
One person stationed at poster; others evaluate
Rubric available
Paper due same time

5. Last In-Class Assignment
Wednesday, December 13th in studio
Available on-line
Read papers before coming to class; bring them with you.
May make any notes you like on the papers
Goal: to evaluate scientific method and data

6. Outline Ed’s demos Balancing Redox Reactions Electrochemistry
Electrochemical cells and Standard Hydrogen Electrodes
Nernst
Quantifying current

7. A
Ed’s Demos

8. Oxidation States of Vanadium: Reduction of V5+ to V+2
Reaction 1
Zn (s) + 2 VO3- (aq) + 8 H3O+ (aq) ↔ 2 VO2+ (aq) + Zn+2 (aq) + 12 H2O (l)
Reaction 2
Zn (s) + 2 VO2+ (aq) + 8 H3O+ (aq) ↔ 2 V3+ (aq) + Zn+2 (aq) + 6 H2O (l)
Reaction 3
Zn (s) + 2 V3+ (aq) ↔ 2 V2+ + Zn+2 (aq)
V+5 (aq) → V+4 (aq) yellow to green
V+4 (aq) → V+3 (aq) green to blue
V+3 (aq) → V+2 (aq) blue to violet

9. Oxidation States of Manganese: Mn+7, Mn+6, Mn+4, and Mn+2
+7 (purple) to +2 (colorless)
2 MnO4- (aq) + H+ (aq) + 5 HSO3- (aq) ↔ 2 Mn+2 (aq) + 5 SO4-2 (aq) + 3 H2O(l)
+ 7 (purple) to +4 (brown)
OH- + 2 MnO4- (aq) + 3 HSO3- (aq) ↔ 2 MnO2 (s) + 3 SO4-2 (aq) + 2 H2O(l)
+ 7 (purple) to + 6 (green)
2 MnO4- (aq) + 3 OH- + HSO3- (aq) ↔ 2 MnO4-2(aq) + SO4-2 (aq) + 2 H2O(l)

10. Thinking back…. What happened when Na(s) was added to water?
Na(s) + H2O(l)  Na+ (aq) H2(g) + OH-(aq)
Determine the oxidation state of each reactant and product
What was oxidized?
What was reduced?

11. Balancing Redox ReactionsM
Balancing Redox Reactions
When KMnO4 (potassium permanganate) is mixed with Na2C2O4 (sodium oxalate) under acidic conditions, Mn+2(aq) ions and CO2(g) form.
The unbalanced chemical equation is:
KMnO4(aq) + Na2C2O4(aq)  Mn+2(aq) + CO2(g) + K+(aq) + Na+(aq)
K+ and Na+ are spectator ions, so we can ignore them at this point.
MnO4- (aq) + C2O4-2(aq)  Mn+2(aq) + CO2(g)

12. Half-Reactions Reduction reaction Oxidation reaction M
MnO4- (aq) + C2O4-2(aq)  Mn+2(aq) + CO2(g)
Reduction reaction
Oxidation reaction
Reduction rxn Mn
Oxidation C

13. Reduction reaction M MnO4-  Mn+2
Step 1: Balance all elements other than oxygen and hydrogen.
Step 2: Balance the oxygens by adding water.
Step 3: Balance the hydrogens using H+
Step 4: Balance the electrons
Mn+7 on reactants side
Mn+2 on products side
Step 5: Check charge balance and elemental balance

14. M
Oxidation reaction
C2O4-2  CO2

15. Combine Half Reactions
5 e H+ + MnO4-  Mn H2O
C2O4-2  2 CO e-
16 H+ + 2MnO4- +5C2O4-2 to 2Mn H2O CO2
We are assuming the reaction takes place under acidic conditions!

16. Balancing in Base 5 e- + 8H+ + MnO4-  Mn+2 + 4 H2O
Change H+ to water by adding OH- to each side
5 e H+ + MnO4-  Mn H2O
C2O4-2  2 CO e-

17. M
NO2-(aq) + Cr2O → Cr+3(aq) + NO3-(aq) acidic soln

18. Electrochemical Cells

19. A
Definitions
Electrochemical cell: A combination of anode, cathode, and other materials arranged so that a product-favored redox reaction can cause a current to flow or an electric current can cause a reactant-favored redox reaction to occur
Voltaic cell (battery): An electrochemical cell or group of cells in which a product-favored redox reaction is used to produce an electric current.
Galvanic cell: A cell in which an irreversible chemical reaction produces electrical current
Electrolytic cell: electrochemical reactions are produced by applying electrical energy

20. A Copper-Zinc battery – What Matters?
Consider reduction potentials:
Cu+2 + 2e- → Cu(s) V
Zn+2 + 2e- → Zn(s) V
Place Zn electrode in copper sulfate solution – What happens?
Copper is plated on Zn electrode
Cu+2 + 2e- → Cu(s) V
Zn(s) → Zn+2 + 2e V
E > 0, spontaneous
Cu+2 + Zn(s) → Zn+2 + Cu(s) 1.1 V
Note, no need for electron to flow external to cell for reaction to occur!!

21. A Copper-Zinc battery – What Matters?
Consider reduction potentials:
Cu+2 + 2e- → Cu(s) V
Zn+2 + 2e- → Zn(s) V
Place Cu electrode in zinc sulfate solution – What happens?
Zn doesn’t plate on copper electrode?!
Cu(s) → Cu+2 + 2e V
Zn+2 + 2e- → Zn(s) V
E < 0, not spontaneous
Zn+2 + Cu(s) → Cu+2 + Zn(s) -1.1 V
No reaction occurs !!

22. A
Fig. 19-3, p.918

23. A
                                                                                           
What are the ½ reactions?
What is the overall reaction?
Identify the oxidation, reduction, anode, and cathode

24. SHE: Standard Hydrogen Electrode
2 H3O+(aq, 1.00 M) + 2e- <->
H2(g, 1 atm) + 2H2O(l)
Eo = 0V
Standard conditions:
1M, 1atm, 25oC
Redox Reactions & Galvanic Cells
In order to compare the potential difference when different metals are used for the anode and the cathode, each half reaction is compared to a standard half reaction.  This process is like choosing par in golf.  The par for a hole in golf is set at the standard.  Golfers scores are then measured relative to par.
The freedom to set the cell potential difference equal to zero for this reaction is the same freedom we have to set any height equal zero wherever is convenient when we work with gravitational potential energy. We have this freedom because we are not interested in the absolute potential of a reaction, rather we are only interested in the magnitude of the change in potential.
Fig. 19-7, p.922

25. Measuring Relative Potentials
                                                                           
Measuring Relative Potentials
Table of Standard
Reduction Potentials

26. Standard Reduction Potentials
What is the standard potential of a Au+3/Au/Mg+2/Mg cell?

27. A
The half-reaction with the more positive standard reduction potential occurs at the cathode as reduction.
The half-reaction with the more negative standard reduction potential occurs at the anode as oxidation.